1. CHE 112 - MODULE 1 CHAPTER 12 LECTURE NOTES

2. Properties of Gases  Gases completely fill their container in a uniform manner  Gases are compressible  Gases have low densities  Gases will exert pressure as the gas molecules collide with the inner surfaces of their container

3. Kinetic Molecular Theory of Gases  A gas is composed of particles that are very small compared with the average distance between the individual particles.  Gas particles are in constant straight line, random motion.  Collisions are elastic.  Attractions between particles are negligible.  The average kinetic energy (energy of motion) of these particles is directly proportional to the temperature in Kelvin.

4. Boyle’s Law  P  1/ V; PV = k when T is constant  P = pressure & V = volume  The pressure is inversely proportional to the volume.  As the pressure increases, the volume decreases; and as the pressure decreases, the volume increases.  P 1 V 1 = P 2 V 2

5. Charles Law  V  T; V = kT when P is constant  V = volume & T = temp in Kelvin  The volume is directly proportional to the temperature.  As the temp increases, the volume increases: and as the temp decreases, the volume decreases.  V 1 /T 1 = V 2 /T 2

6. Combined Gas Law  Combination of Boyle’s and Charle’s Laws  P 1 V 1 / T 1 = P 2 V 2 / T 2  Used when all the conditions of the same gas changes, and nothing remains constant.

7. Avogadro’s Law  The molar volume is defined as the volume of 1 mole = 22.4 L for any gas at Standard Temperature and Pressure (STP)  STP means 273.15K (0°C) and 1 atm (760mm Hg)

8. Ideal Gas Law PV = nRT  P = pressure (atm)  V = volume (L)  n = number of moles (mol)  R = gas constant = 0.08206 (L·atm/mol·K)  T = temperature (K)

9. Real Gas Law  As temperatures and pressures deviate from STP, the Ideal Gas Law no longer holds true.  The Van der Waals equation is used to compensate for intermolecular forces and molecular volume in non-ideal conditions. (P+a[n/V] 2 )(V-bn) = nRT

10. CHE 112 - MODULE 1 CHAPTER 13 LECTURE NOTES

11. Intermolecular Forces  Dipole-dipole interactions –polar molecules attract by charge  Dipole-induced dipole forces –a dipole can be induced by a polar species  London dispersion forces –momentary attraction of non-polar molecules  Hydrogen bonding –attraction of H-O, H-N and H-F with other highly electronegative atoms

12. Water and Hydrogen Bonding  The O of a water molecule attaches to four other water molecules  The hydrogen bonded water creates a tetrahedral of H around the O  Water freezes into a water lattice that is ordered and has spaces within the repeated lattice crystal

13. DNA and Hydrogen Bonding  Deoxyribonucleic acid (DNA) is a chain of phosphates linked to sugars  Bonded to each sugar is either a thymine, guanine, cytosine, or adenine base molecule  The base molecules on one chain interact through hydrogen bonding with base molecules on another chain  This pairing of chains produces the double helix of DNA

14. Properties of Liquids  Vaporization =  H° vap standard molar heat of vaporization –liquid  gas, endothermic process –liquid  vapor, vapor pressure –boiling point  Surface tension - energy to break the surface  Capillary action - adhesive forces  Viscosity - resistance to flow

15. Vapor Pressure  Vapor pressure = the partial pressure of the vapor over the liquid, measured at equilibrium –As surface molecules gain sufficient KE, they escape the liquid into vapor above the liquid –As the # of vapor molecules increase, more collisions will occur above the liquid –More molecules will collide at the surface and condense back into the liquid state –Rate of condensation = rate of vaporization –At this point, the partial pressure exerted = vapor pressure of the liquid –As the temperature , KE , vapor pressure 

16. Boiling Point  Boiling point = the temperature at which the vapor pressure of a liquid = the vapor pressure exerted on the liquid* –As temperature , vapor pressure  ; when the vapor pressure = atmospheric pressure, stable bubbles of vapor form –The liquid begins to “boil” –Once boiling begins, it will continue to boil as long as heat is supplied or until all the liquid has been converted to vapor *At 1atm, water will boil at 100°C, but at the top of Whiteface, water will boil at a lower temperature

17. Types of Solids  Ionic compounds  Metallic materials  Molecular compounds  Network solids  Amorphous materials

18. Crystalline vs Amorphous  Crystalline = has a defined structure  Amorphous = random (undefined structure)

19. Cubic Units  Simple cube  Face-centered cube  Body centered cube

20. Phase Diagrams  Phase Diagram shows all the phases of matter and their relationships for different types of matter –Solid –Liquid –Gas –Triple point –Critical point

21. Changes of States  Melting (solid > liquid)  Freezing (liquid > solid)  Vaporization (liquid > vapor)  Condensation (vapor > liquid)  Sublimation (solid > vapor)  Condensation by deposition (vapor > solid)

22. CHE 112 - MODULE 1 CHAPTER 14 LECTURE NOTES

23. Units of Concentration  Molarity (M) = moles of solute Liters of solution  Molality (m) = moles of solute Kg of solvent  Mole fraction (X A ) = moles of A moles of solution

24. Units of Concentration (cont.)  Weight % A = mass of A X 100% total mass  Parts per million (ppm) = 1g of matter mass million g 1ppm = 1mg/L = 1  g/ml

25. Solutions  Saturated - solution containing the maximum amount of dissolved solvent  Unsaturated - Less than saturated  Supersaturated - more than saturated

26. Liquid/Liquid Interaction  Miscible - liquids that mix well due to like polarity  Immiscible - liquids that do not mix because one is polar and the other is non-polar “Likes dissolve likes”

27. Solids Dissolved in Liquids  Hydrated - ionic solids dissolve when surrounded by polar water molecules in solution  Heat of Solution (  H soln ) = net energy change for the solution process –Exothermic rxn yield more soluble cmp –Endothermic rxn yield insoluble cmp

28. Enthalpy of Hydration  The energy involved with hydrating an ion depends on 3 things: –distance between the ion & dipole (closer the stronger attraction) –charge on the ion (higher the stronger) –polarity of the solvating molecule (greater the magnitude of the dipole the stronger)

29. Vapor Pressure Lowering  VPL = a colligative property (concentration dependent) equal to the vapor pressure of pure solvent minus the vapor pressure of the solution –Occurs when a nonvolatile solute is mixed with a solvent that have similar structure –  P solv = P solv -  P solv

30. Raoult’s Law  P solv = X solv P  solv – P solv  X solv –Vapor pressure of the solution is lower than the that of the pure solvent –CD-ROM Screen 14.7   P solv = -X solv P  solv

31. Freezing Point Depression  Freezing point depression is a colligitive property equal to the freezing point of the pure solvent minus the freezing point of the solution   T fp = k fp m solute –  T fp = freezing point depression –k fp = freezing point depression constant (for water it is -1.86°C/m) –m solute = molal concentration

32. Boiling Point Elevation  Boiling point elevation is a colligitive property equal to the boiling point of the solution minus the boiling point of the pure solvent   T bp = k bp m solute –  T bp = boiling point elevation –k bp = boiling point elevation constant (for water it is 0.512°C/m) –m solute = molal concentration

33. Colloids  Dispersion of particles of one substance (dispersed phase) throughout another substance or solution (the continuous phase)  Appears homogeneous  Particle sizes are between 10 – 2000 angstroms (1x10 -10 m) in diameter

34. Micelle  Aggregate of molecules that have a polar end and a nonpolar end. –Nonpolar ends will meet together in the middle, while the polar ends will be attracted outward toward the water molecules